Idea Transcript
Writing Electron Configurations
The arrangement of elements in the periodic table reflects the arrangement of electrons in an atom. Each period begins with an atom that has an electron in a new energy level and with the exception of the first period, each period ends with an atom that has a filled set of p orbitals.
To write the electron configuration of an element, you must fill the sublevels in order of increasing energy. If you follow the arrows in either of the two types of mnemonics shown below, you will get correct configurations for most elements.
1s
1s 2s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6f
6s
6p
6d
7d
7s
7p
2p 3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
7p
You also need to know how many orbitals are in each sublevel and that each orbital can contain two electrons of opposite
Answers 1. a. 1s22s22p63s23p1 b. 1s22s22p6 c. 1s22s22p63s23p63d 104s24p6 4d 105s25p2 d. 1s22s22p63s23p64s1 2. a. [Ne]3s23p2 b. [Kr]5s1 c. [Kr]4d 105s25p3 d. [Ar]3d 104s24p3
spins. As shown in the following table, the sublevels s, p, d, and f have 1, 3, 5, and 7 available orbitals, respectively.
SUBLEVEL
s
p
No. of orbitals
1
3
5
7
No. of electrons
2
6
10
14
d
f
Sample Problem Write the full electron configuration for phosphorus.
]
]
]
]
The atomic number of phosphorus is 15, so a phosphorus atom has 15 protons and electrons. Assign each of the 15 electrons to the appropriate sublevels. The final sublevel can be unfilled and will contain the number of valence electrons. 6 3 1s2 2s2 2p 3s2 3p
]
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Math Tutor
M at h T u t o r
2e- + 2e- + 6e- + 2e- + 3e- = 15eSo, the full electron configuration of phosphorus is 1s22s22p63s23p3.
Answers in Appendix E
1. Write full electron configurations for the following elements. a. aluminum c. tin b. neon d. potassium
2. Use noble gas symbols to write shorthand electron configurations for the following elements. a. silicon c. antimony b. rubidium d. arsenic
Math Tutor
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The Periodic Law 157
C h a p t e r s u mm a ry
CHAPTER 5 SECTION 1
Summary
PREMIUM CONTENT
Interactive Review HMDScience.com
History of the Periodic Table
• The periodic law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers. • The periodic table is an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column.
Review Games Concept Maps
KEY TERMS
periodic law periodic table lanthanide actinide
• The columns in the periodic table are referred to as groups.
SECTION 2
Electron Configuration and the Periodic Table
• The rows in the periodic table are called periods. • Many chemical properties of the elements can be explained by the configurations of the elements’ outermost electrons. • The noble gases exhibit unique chemical stability because their highest occupied levels have an octet of electrons, ns2np6 (with the exception of helium, whose stability arises from its highest occupied level being completely filled with two electrons, 1s2).
KEY TERMS
alkali metals alkaline-earth metals transition elements main-group elements halogens
• Based on the electron configurations of the elements, the periodic table can be divided into four blocks: the s-block, the p-block, the d-block, and the f-block.
SECTION 3
Electron Configuration and Periodic Properties
• The groups and periods of the periodic table display general trends in the following properties of the elements: electron affinity, electronegativity, ionization energy, atomic radius, and ionic radius. • The electrons in an atom that are available to be lost, gained, or shared in the formation of chemical compounds are referred to as valence electrons. • In determining the electron configuration of an ion, the order in which electrons are removed from the atom is the reverse of the order given by the atom’s electron-configuration notation.
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1. a. Cannizzaro developed a standard method for measuring atomic masses, which allowed chemists to search for periodic trends among elements. b. Mendeleev organized elements according to increasing atomic mass and noticed that similar properties appeared periodically. c. Moseley discovered that nuclear charge (i.e., atomic number), not atomic mass, should be the basis for organizing the periodic table. 158 Chapter 5
atomic radius ion ionization ionization energy electron affinity cation anion valence electron electronegativity
Chapter 5
Review Answers Untitled-46 158
KEY TERMS
2. The physical and chemical properties of the elements are periodic functions of their atomic numbers. 3. Groups of elements exhibit similar chemical and physical properties and behavior. 4. a. Generally, the configurations of the outermost electron shells of elements within the same group are the same. There are a number of exceptions to this rule, however, among the transition elements. b. Their outer shells are completely filled.
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CHAPTER 5
C HAPTER RE V I E W
Review
SECTION 1
History of the Periodic Table REVIEWING MAIN IDEAS 1. Describe the contributions made by the following scientists to the development of the periodic table: a. Stanislao Cannizzaro b. Dmitri Mendeleev c. Henry Moseley 2. State the periodic law. 3. How is the periodic law demonstrated within the groups of the periodic table? SECTION 2
Electron Configuration and the Periodic Table REVIEWING MAIN IDEAS 4. a. How do the electron configurations within the same group of elements compare? b. Why are the noble gases relatively unreactive? 5. What determines the length of each period in the periodic table? 6. What is the relationship between the electron configuration of an element and the period in which that element appears in the periodic table? 7. a. What information is provided by the specific block location of an element? b. Identify, by number, the groups located within each of the four block areas. 8. a. Which elements are designated as the alkali metals? b. List four of their characteristic properties. 9. a. Which elements are designated as the alkalineearth metals? b. How do their characteristic properties compare with those of the alkali metals? 10. a. Write the group configuration notation for each d-block group. b. How do the group numbers of those groups relate to the number of outer s and d electrons?
11. What name is sometimes used to refer to the entire set of d-block elements? 12. a. What types of elements make up the p-block? b. How do the properties of the p-block metals compare with those of the metals in the s- and d-blocks? 13. a. Which elements are designated as the halogens? b. List three of their characteristic properties. 14. a. Which elements are metalloids? b. Describe their characteristic properties. 15. Which elements make up the f-block in the periodic table? 16. a. What are the main-group elements? b. What trends can be observed across the various periods within the main-group elements?
PRACTICE PROBLEMS 17. Write the noble-gas notation for the electron configuration of each of the following elements, and indicate the period in which each belongs. a. Li c. Cu e. Sn b. O d. Br 18. Without looking at the periodic table, identify the period, block, and group in which the elements with the following electron configurations are located. (Hint: See Sample Problem A.) a. [Ne]3s23p4 b. [Kr]4d105s25p2 c. [Xe]4f 145d106s26p5 19. Based on the information given below, give the group, period, block, and identity of each element described. (Hint: See Sample Problem B.) a. [He]2s2 b. [Ne]3s1 c. [Kr]5s2 d. [Ar]4s2 e. [Ar]3d54s1 20. Without looking at the periodic table, write the expected outer electron configuration for each of the following elements. (Hint: See Sample Problem C.) a. Group 7, fourth period b. Group 3, fifth period c. Group 12, sixth period Chapter Review
5. The length of a period is determined by the total number of electrons that can fill the outer sublevels of the elements of that period. 6. An element’s period corresponds to its highest occupied main energy level. 7. a. the type of sublevel being filled in successive elements of that block b. s-block: Groups 1 and 2; p-block: Groups 13–18 (except for He); d-block: Groups 3–12 (except for the f‑block elements); f-block: those elements of the sixth and seventh periods between Groups 3 and 4
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8. a. the Group 1 elements b. The Group 1 elements are extremely 5/9/2011 10:55:29 AM reactive. They react vigorously with water, they are silvery in appearance, and each is soft enough to be cut with a knife. 9. a. the Group 2 elements b. The Group 2 elements are harder, denser, and stronger than the Group 1 elements. They also have higher melting points. Group 2 elements are less reactive than Group 1 elements.
10. a. Group 3: (n – 1)d 1ns2 Group 4: (n – 1)d2ns2 Group 5: (n – 1)d3ns2 Group 6: (n – 1)d5ns1 for Cr and Mo (n – 1)d4ns2 for W and Sg Group 7: (n – 1)d5ns2 Group 8: (n – 1)d6ns2 Group 9: (n – 1)d 7ns2 Group 10: (n – 1)d8ns2 Group 11: (n – 1)d 10ns1 Group 12: (n – 1)d 10ns2 b. The group number equals the sum of the outer s and d electrons. 11. the transition elements 12. a. The p-block consists of nonmetals at the right, metalloids in the middle, and metals at the left. b. The p‑block metals are generally harder and more dense than the s‑block metals but softer and less dense than the d‑block metals. 13. a. the Group 17 elements b. The halogens are the most-reactive nonmetals; they react vigorously with most metals to form salts. They are also among the most electronegative elements. 14. a. B, Si, Ge, As, Sb, Te b. The metalloids are mostly brittle solids with electrical conductivities intermediate between those of metals (good conductors) and nonmetals (nonconductors). 15. the lanthanides and the actinides 16. a. the elements of the s and p-blocks (plus hydrogen and helium) b. decrease in atomic size, increase in ionization energy, increase in electron affinity, decrease in cationic size, decrease in anionic size, increase in electronegativity 17. a. [He]2s1, second period b. [He]2s22p4, second period c. [Ar]3d 104s1, fourth period d. [Ar]3d 104s24p5, fourth period e. [Kr]4d 105s25p2, fifth period 18. a. third period, p-block, Group 16 b. fifth period, p-block, Group 14 c. sixth period, p-block, Group 17 19. a. Group 2, second period, s-block, Be b. Group 1, third period, s-block, Na c. Group 2, fifth period, s-block, Sr d. Group 2, fourth period, s-block, Ca e. Group 6, fourth period, d-block, Cr 20. a. 3d54s2, b. 4d15s2, c. 4f145d106s2 The Periodic Law 159
c h a p t e r r e vi e w 21. a. p-block, third period, Group 13, Al, metal, high reactivity b. p-block, fourth period, Group 18, noble gases, Kr, nonmetal, low reactivity c. d-block, fifth period, Group 11, Ag, metal, low reactivity d. f-block, sixth period, between Groups 3 and 4, Ce, metal, high reactivity 22. a. one half the distance between the nuclei of two bonded identical atoms b. They decrease. c. As electrons are added to s and p sublevels in the same main energy level, the increasing positive charge of the nucleus pulls electrons closer to the nucleus, resulting in decreasing atomic radii. 23. a. They generally increase. b. Down a group, the outer electrons of each element occupy comparable sublevels in successively higher main energy levels farther from the nucleus. 24. a. a charged atom or a charged group of bonded atoms b. any process that results in the formation of an ion c. the energy required to remove one electron from a neutral atom of an element d. the energy required to remove an electron from a 1+ ion 25. a. They increase across a period and decrease down a group. b. Across a period, the increasing nuclear charge attracts electrons in the same energy level more strongly and makes them more difficult to remove. Down a group, the electrons to be removed from each successive element are farther from the nucleus in increasingly higher energy levels and are thus more easily removed. 26. a. the energy taken in or given off when an electron is added to an atom b. Electron affinity values are either negative or positive. A negative sign indicates that energy is given off; a positive sign indicates that energy is taken in. 27. a. A cation is a positive ion, and an anion is a negative ion.
160 Chapter 5
CHAPTER REVIEW 21. Identify the block, period, group, group name (where appropriate), element name, element type, and relative reactivity for the elements with the following electron configurations. (Hint: See Sample Problem D.) a. [Ne]3s23p1 b. [Ar]3d104s24p6 c. [Kr]4d105s1 d. [Xe]4f 15d16s2 SECTION 3
Electron Configuration and Periodic Properties
29. For each of the following groups, indicate whether electrons are more likely to be lost or gained in compound formation, and give the number of such electrons typically involved. a. Group 1 d. Group 16 b. Group 2 e. Group 17 c. Group 13 f. Group 18 30. a. What is electronegativity? b. Why is fluorine special in terms of electronegativity? 31. Identify the most- and least-electronegative groups of elements in the periodic table.
REVIEWING MAIN IDEAS
PRACTICE PROBLEMS
22. a. What is meant by atomic radius? b. What trend is observed among the atomic radii of main-group elements across a period? c. Explain this trend.
32. Of cesium, Cs, hafnium, Hf, and gold, Au, which element has the smallest atomic radius? Explain your answer in terms of trends in the periodic table. (Hint: See Sample Problem E.)
23. a. What trend is observed among the atomic radii of main-group elements down a group? b. Explain this trend.
33. a. Distinguish between the first, second, and third ionization energies of an atom. b. How do the values of successive ionization energies compare? c. Why does this occur?
24. Define each of the following terms: a. ion b. ionization c. first ionization energy d. second ionization energy
34. Without looking at the electron affinity table, arrange the following elements in order of decreasing electron affinities: C, O, Li, Na, Rb, and F.
25. a. How do the first ionization energies of main-group elements vary across a period and down a group? b. Explain the basis for each trend. 26. a. What is electron affinity? b. What signs are associated with electron affinity values, and what is the significance of each sign? 27. a. Distinguish between a cation and an anion. b. How does the size of each compare with the size of the neutral atom from which it is formed? 28. a. What are valence electrons? b. Where are such electrons located?
35. a. Without looking at the ionization energy table, arrange the following elements in order of decreasing first ionization energies: Li, O, C, K, Ne, and F. b. Which of the elements listed in (a) would you expect to have the highest second ionization energy? Why? 36. a. Which of the following cations is least likely to form: Sr2+, Al3+, K2+? b. Which of the following anions is least likely to form: I-, Cl-, O2-? 37. Which element is the most electronegative among C, N, O, Br, and S? Which group does it belong to? (Hint: See Sample Problem G.) 38. The two ions K+ and Ca2+ each have 18 electrons surrounding the nucleus. Which would you expect to have the smaller radius? Why?
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b. Cations are always smaller than the atoms from which they are formed; anions are Untitled-46 160 always larger. 28. a. Valence electrons are atomic electrons available to be lost, gained, or shared in the formation of chemical compounds. b. Valence electrons are located in an atom’s outermost energy level. 29. a. lost, 1 d. gained, 2 b. lost, 2 e. gained, 1 c. lost, 3 f. neither lost nor gained, 0 30. a. Electronegativity is the ability of an atom in a chemical compound to attract
electrons from other atoms. b. Fluorine is the most electronegative 5/9/2011 element and is arbitrarily assigned an electronegativity of 4.0. The values for all other elements are assigned in relation to this value. 31. Most-electronegative: Group 17, least-electronegative: Group 1 32. Gold; atomic radii decrease across a period, and gold is farthest to the right in the sixth period, in which all 3 elements are found. 33. a. First ionization energy is the energy required to remove one electron from an
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CHAPTER REVIEW
REVIEWING MAIN IDEAS 39. Without looking at the periodic table, identify the period, block, and group in which each of the following elements is located. a. [Rn]7s1 b. [Ar]3d 24s 2 c. [Kr]4d105s1 d. [Xe]4f 145d 96s1 40. a. Which elements are designated as the noble gases? b. What is the most significant property of these elements?
47. Why do the halogens readily form 1–ions? 48. Identify which trends in the diagrams below describe atomic radius, ionization energy, electron affinity, and electronegativity. Increases
Increases
Mixed Review
a.
41. Which of the following does not have a noble-gas configuration: Na+, Rb+, O2-, Br- Ca+, Al3+, S2-?
44. Use the periodic table to describe the chemical properties of the following elements: a. fluorine, F b. xenon, Xe c. sodium, Na d. gold, Au 45. For each element listed below, determine the charge of the ion that is most likely to be formed and the identity of the noble gas whose electron configuration is thus achieved. a. Li e. Mg i. Br b. Rb f. Al j. Ba c. O g. P d. F h. S
Increases
43. Write the noble-gas notation for the electron configuration of each of the following elements, and indicate the period and group in which each belongs. a. Mg b. P c. Sc d. Y
Increases
b.
Increases
Increases
42. a. How many groups are in the periodic table? b. How many periods are in the periodic table? c. Which two blocks of the periodic table make up the main-group elements?
c.
49. The electron configuration of argon differs from those of chlorine and potassium by one electron each. Compare the reactivity of these three elements.
46. Describe some differences between the s-block metals and the d-block metals.
Chapter Review
atom, which is neutral. Second ionization energy is the energy required to remove an electron from a 1+ ion. Third ionization energy is the energy required to remove an electron from a 2+ ion. b. Each successive ionization energy is larger than the preceding one. c. Each successive electron must be removed from a particle with a greater positive charge. 34. In order of decreasing electron affinities, the elements are F, O, C, Li, Na, and Rb. 35. a. Ne, F, O, C, Li, K
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b. Li and K would have the highest second ionization energies, because in5/9/2011 both cases 10:55:32 AM the second electron must come from a completely filled noble-gas electron configuration. Of the two, Li would have the higher second ionization energy, because the Li+ ion is smaller than the K+ ion. 36. a. K2+, b. O2‑ 37. O; Group 16 38. Ca2+; because of the greater attraction that its 20 protons exert on the 18 electrons 39. a. seventh period, s-block, Group 1
C HAPTER RE V I E W b. fourth period, d-block, Group 4 c. fifth period, d-block, Group 11 d. sixth period, d-block, Group 10 40. a. the Group 18 elements b. They are very unreactive. 41. Ca+ 42. a. 18; b. 7; c. the s and p-blocks 43. a. [Ne]3s2, third period, Group 2 b. [Ne]3s23p3, third period, Group 15 c. [Ar]3d 14s2, fourth period, Group 3 d. [Kr]4d 15s2, fifth period, Group 3 44. a. It is a highly reactive nonmetal in Group 17, Period 2. It needs one electron to achieve a noble gas configuration, which explains its high electron affinity. It has a high ionization energy, and thus the formation of positive ions is not likely. Fluorine is the most electronegative element. b. It is a nonmetal of low reactivity in Group 18, Period 5. It has a filled outer energy level, so there is little tendency to lose, gain, or share electrons. It has the highest ionization energy of the elements in the fifth period. c. It is a highly reactive metal in Group 1, Period 3. It has a low first ionization energy, because losing an electron to form Na+ gives it a noble gas configuration. It has a low electron affinity and a low electronegativity. d. It is a metal of fairly low reactivity in Group 11, Period 6. Its ionization energy is relatively low, so it forms positive ions. 45. a. 1+, He f. 3+, Ne b. 1+, Kr g. 3‑, Ar c. 2‑, Ne h. 2‑, Ar d. 1‑, Ne i. 1‑, Kr e. 2+, Ne j. 2+, Xe 46. The d‑block metals are harder, denser, and less reactive, and except for mercury, they have higher melting points than s‑block metals do. 47. In doing so, they achieve a noble-gas configuration. 48. ionization energy, electron affinity, electronegativity: diagram a; atomic radius: diagram b 49. Chlorine and potassium are reactive elements, whereas argon is unreactive. Chlorine readily gains one electron to achieve the configuration of argon, and potassium readily loses one electron to achieve the argon configuration. The Periodic Law 161
c h a p t e r r e vi e w
CHAPTER REVIEW CRITICAL THINKING
50.
13 β
9 Λ
4 ∆
12 __ γ
23 ξ
51. 5 52. 8 53. From left to right, the elements are organized according to their atomic numbers. Elements are organized into families or columns according to their properties. The elements exhibit increasingly nonmetallic behavior from left to right across the table. Their physical states suggest that melting points and boiling points also increase from left to right. 54. Na is a liquid at 97.8°C. K is a liquid at 63.25°C. Rb is a liquid at 38.89°C. Cs is a liquid at 28.5°C. Fr is a liquid at 27°C. Hg is a liquid at ‑38.8°C. Ga is a liquid at 29.77°C. N as N2 is a gas at ‑195.8°C. P is a liquid at 44.1°C. O as O2 is a gas at ‑182.962°C. F as F2 is a gas at ‑188.14°C. Cl as Cl2 is a gas at ‑34.6°C. Br as Br2 is a gas at 58.78°C. 55. a. Cr3+, violet or green; Cr6+, yellow or orange b. rubies c. copper, blue; cadmium, yellow; cobalt, pink; zinc, white; nickel, green d. Gold, silver, platinum, palladium, iridium, rhodium, ruthenium, osmium; they have low reactivity. 56. Possible responses include Friedrich Dorn’s discovery of radon in 1900, Moseley’s determination of the nuclear charge (i.e., atomic number) of each element by 1913, the discovery of the lanthanides by 1913, the synthesis in 1940 of the first element beyond uranium (neptunium) in the periodic table by Philip Abelson and Edwin McMillan, the creation of plutonium in 1941 by Seaborg, and the synthesis of element 109 in Germany in 1982.
162 Chapter 5
USING THE HANDBOOK
As a member on the newly-inhabited space station Alpha, you are given the task of organizing information on newly discovered elements as it comes in from the laboratory. To date, five elements have been discovered and have been assigned names and symbols from the Greek alphabet. An analysis of the new elements has yielded the following data:
Element Atomic Atomic Properties name no. mass Epsilon ε
23
47.33
nonmetal, very reactive, produces a salt when combined with a metal, gaseous state
Beta β
13
27.01
metal, very reactive, soft solid, low melting point
Gamma γ
12
25.35
nonmetal, gaseous element, extremely unreactive
Delta ∆
4
7.98
nonmetal, very abundant, forms compounds with most other elements
Lambda Λ
9
16.17
metal, solid state, good conductor, high luster, hard and dense
54. Review the boiling point and melting point data in the tables of the Elements Handbook (Appendix A). Make a list of the elements that exist as liquids or gases at the boiling point of water, 100°C. 55. Because transition metals have vacant d orbitals, they form a greater variety of colored compounds than do the metals of Groups 1 and 2. Review the section of the Elements Handbook (Appendix A) on transition metals, and answer the following: a. What colors are exhibited by chromium in its common oxidation states? b. What gems contain chromium impurities? c. What colors are often associated with the following metal ions: copper, cadmium, cobalt, zinc, and nickel? d. What transition elements are considered noble metals? What are the characteristics of a noble metal?
RESEARCH AND WRITING 56. Prepare a report tracing the evolution of the current periodic table since 1900. Cite the chemists involved and their major contributions. 57. Write a report describing the contributions of Glenn Seaborg toward the discovery of many of the actinide elements.
ALTERNATIVE ASSESSMENT 50. Applying Models Create a periodic table based on the properties of the five new elements. 51. Predicting Outcomes Using your newly created periodic table, predict the atomic number of an element with an atomic mass of 11.29 that has nonmetallic properties and is very reactive.
58. Construct your own periodic table or obtain a poster that shows related objects, such as fruits or vegetables, in periodic arrangement. Describe the organization of the table and the trends it illustrates. Use this table to make predictions about your subject matter.
52. Predicting Outcomes Predict the atomic number of an element having an atomic mass of 15.02 that exhibits metallic properties but is softer than lambda and harder than beta. 53. Analyzing Information Analyze your periodic table for trends, and describe those trends.
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57. Reports should mention Seaborg’s various roles in the syntheses of curium (element 96) Untitled-46 162 in 1944, californium (element 98) in 1950, mendelevium (element 101) in 1955, and nobelium (element 102) in 1958. 58. Student examples may vary, but should include isolated trends arranged in a periodic manner—by columns and rows, as in the periodic table of the elements.
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